Intermolecular Forces, Liquids & Solids

Intramolecular vs Intermolecular Forces

Intramolecular forces are the covalent, ionic, and metallic bonds that hold atoms together within a molecule or lattice; these are strong (100s of kJ/mol) and determine chemical identity. Intermolecular forces (IMFs) are the much weaker attractions (a few to tens of kJ/mol) between separate molecules or ions that determine physical properties like melting point, boiling point, and solubility. Melting or boiling a substance breaks IMFs, not covalent bonds — this is why water boils at 100 °C while breaking O–H bonds requires far more energy. The stronger the net IMFs, the higher the melting point, boiling point, viscosity, and surface tension, and the lower the vapor pressure.

• •Intramolecular = within the molecule (ionic, covalent, metallic bonds)
• •Intermolecular = between molecules (dispersion, dipole-dipole, H-bonding, ion-dipole)
• •Phase changes only break/reform IMFs, never covalent bonds
• •Stronger IMFs → higher mp/bp, higher viscosity, higher surface tension, lower vapor pressure

Dispersion (London) Forces and Polarizability

Dispersion forces arise from instantaneous, temporary dipoles created by the constant motion of electrons; a momentary imbalance in one molecule induces a complementary dipole in a neighbor. These forces exist in all molecules and atoms, polar or nonpolar, and are the only IMF present in nonpolar species like noble gases and hydrocarbons. Polarizability — how easily an electron cloud is distorted — increases with larger atomic/molecular size (more electrons, more diffuse cloud) and with molar mass, so dispersion forces strengthen down a family or with increasing molar mass. Molecular shape matters too: elongated, more contact-area molecules (like n-pentane) have stronger dispersion forces and higher boiling points than compact, branched isomers (like neopentane) of the same molar mass.

• •Present in ALL molecules/atoms; only IMF in nonpolar species
• •Strength increases with molar mass and number of electrons (higher polarizability)
• •Elongated shapes have more surface contact → stronger dispersion than compact/branched isomers
• •Explains why F₂ is a gas but I₂ is a solid at room temperature

Dipole–Dipole Forces and Hydrogen Bonding

Polar molecules possess a permanent dipole, and dipole–dipole forces are attractions between the positive end of one polar molecule and the negative end of another; these add to dispersion forces and generally raise boiling point relative to a nonpolar molecule of similar mass. Hydrogen bonding is a special, unusually strong dipole-dipole interaction that occurs only when H is covalently bonded directly to N, O, or F, because these small, highly electronegative atoms create a very concentrated positive charge on the tiny hydrogen. Hydrogen bonding explains why H₂O, NH₃, and HF have anomalously high boiling points compared to other hydrides in their groups (e.g., H₂O vs H₂S), and it is responsible for water's high surface tension, high specific heat, and the expansion of ice upon freezing.

• •Dipole-dipole requires permanent molecular polarity
• •H-bonding requires H bonded to N, O, or F specifically
• •H-bonding is the strongest common IMF (still ~5–10% the strength of a covalent bond)
• •Explains anomalous bp trends: H₂O, NH₃, HF break the pattern set by their heavier group members

Ion–Dipole Forces and Ranking IMFs

Ion–dipole forces occur when an ion (from a dissolved ionic compound) attracts the appropriately charged end of a polar solvent molecule, such as Na⁺ surrounded by the oxygen ends of water molecules; these are the strongest common IMF and explain why ionic compounds dissolve in polar solvents. The overall order of typical IMF strength is dispersion < dipole-dipole < hydrogen bonding < ion-dipole, though a very large, polarizable nonpolar molecule can have stronger net dispersion forces than the dipole-dipole forces of a small polar one. To rank compounds by boiling point or vapor pressure, first identify the strongest IMF present in each, then compare molar mass/polarizability within the same IMF category.

• •Ion-dipole occurs in solutions of ionic compounds in polar solvents (hydration of ions)
• •General strength order: dispersion < dipole-dipole < H-bonding < ion-dipole
• •Always identify the STRONGEST IMF present first when ranking properties
• •Large nonpolar molecules can out-rank small polar ones due to dominant dispersion forces

Liquid Properties: Viscosity, Surface Tension, Capillary Action

Viscosity is a liquid's resistance to flow and increases with stronger IMFs and with molecular shape that allows entanglement (long chains); it decreases with increasing temperature as molecules gain kinetic energy to overcome IMFs. Surface tension is the energy required to increase a liquid's surface area, arising because surface molecules have unbalanced cohesive forces pulling them inward; stronger IMFs give higher surface tension, which is why water beads up and can support a paperclip. Capillary action is the rise of liquid in a narrow tube, driven by the competition between cohesive forces (liquid-liquid) and adhesive forces (liquid-container); water in a glass tube shows a concave (upward-curving) meniscus because adhesion to glass exceeds cohesion, while mercury shows a convex meniscus because cohesion dominates.

• •Viscosity ↑ with stronger IMFs and chain length; ↓ with temperature
• •Surface tension ↑ with stronger IMFs (water is unusually high due to H-bonding)
• •Cohesion = liquid-liquid attraction; adhesion = liquid-surface attraction
• •Concave meniscus: adhesion > cohesion (water/glass); convex meniscus: cohesion > adhesion (mercury/glass)

Vaporization, Vapor Pressure, and Boiling

Vaporization is the escape of high-kinetic-energy molecules from a liquid surface into the gas phase; in a closed container, vapor molecules also condense back, and dynamic equilibrium is reached when the rates of vaporization and condensation are equal. The equilibrium vapor pressure at that point depends only on temperature and the identity of the liquid (weaker IMFs → higher vapor pressure), not on the amount of liquid present. A liquid boils when its vapor pressure equals the external (atmospheric) pressure; the normal boiling point is specifically the temperature at which vapor pressure equals exactly 1 atm, so boiling point drops at high altitude (lower atmospheric pressure) and rises in a pressure cooker.

• •Dynamic equilibrium: rate of vaporization = rate of condensation
• •Vapor pressure depends on temperature and IMF strength, not on liquid volume
• •Boiling occurs when vapor pressure = external pressure
• •Normal boiling point = temperature where vapor pressure = 1 atm exactly

Heating/Cooling Curves and Enthalpies of Phase Change

A heating curve plots temperature versus heat added at constant pressure; sloped segments represent heating within one phase (q = mCΔT) while flat plateaus represent phase changes occurring at constant temperature (q = nΔH_fus or q = nΔH_vap). During a plateau, all added energy goes into overcoming IMFs to separate molecules rather than increasing kinetic energy/temperature, which is why temperature stays constant while ice melts or water boils. Heat of vaporization is always considerably larger than heat of fusion for the same substance because vaporization requires completely separating molecules against all IMFs, whereas melting only needs to disrupt the rigid, ordered lattice enough to allow flow.

• •Sloped segments: q = mCΔT (single phase warming/cooling)
• •Flat plateaus: q = nΔH (phase change at constant T)
• •ΔH_vap > ΔH_fus always, for the same substance
• •Multi-step problems require adding heat contributions from every segment crossed

Phase Diagrams

A phase diagram plots pressure (y-axis) versus temperature (x-axis) and is divided into three regions (solid, liquid, gas) separated by three curves: the solid-liquid (melting/freezing) line, the liquid-gas (vaporization/condensation) line, and the solid-gas (sublimation/deposition) line. The triple point is the unique temperature and pressure at which all three phases coexist in equilibrium simultaneously; the critical point is the temperature and pressure beyond which liquid and gas become indistinguishable, forming a supercritical fluid with properties of both. For most substances the solid-liquid line slopes slightly to the right (solid is denser than liquid, so increasing pressure favors the solid), but water's solid-liquid line slopes to the left (negative slope) because ice is less dense than liquid water, so increasing pressure on ice favors melting it.

• •Three regions (solid, liquid, gas) separated by three boundary curves
• •Triple point: all three phases coexist at one specific T and P
• •Critical point: endpoint of the liquid-gas line; beyond it, supercritical fluid forms
• •Water's negative solid-liquid slope reflects that ice floats (solid is less dense than liquid)

Classifying Solids and Unit Cells

Solids are classified by the particles present and forces holding them: ionic solids (ions held by electrostatic attraction, e.g., NaCl) are hard, brittle, high-melting, and conduct only when molten/dissolved; metallic solids (cations in a sea of delocalized electrons) are malleable, ductile, and conduct electricity as solids; covalent-network solids (atoms linked by a continuous covalent bond network, e.g., diamond, SiO₂) are extremely hard and high-melting with poor conductivity; molecular solids (discrete molecules held by IMFs, e.g., ice, sucrose) are soft and low-melting. Crystalline solids often pack in repeating unit cells: simple cubic has 1 atom/cell (8 corners × 1/8) with coordination number 6; body-centered cubic (BCC) has 2 atoms/cell (8×1/8 + 1 center) with coordination number 8; face-centered cubic (FCC) has 4 atoms/cell (8×1/8 + 6×1/2) with coordination number 12.

• •Ionic: hard, brittle, high mp, conducts only molten/aqueous
• •Metallic: malleable, ductile, conducts as a solid, lustrous
• •Covalent-network: extremely hard, very high mp, poor conductor (diamond, quartz)
• •Molecular: soft, low mp, held together by IMFs only
• •Simple cubic: 1 atom/cell, CN=6; BCC: 2 atoms/cell, CN=8; FCC: 4 atoms/cell, CN=12

Solutions: Like Dissolves Like and Solubility Trends

'Like dissolves like' summarizes that solutes dissolve best in solvents with similar IMFs — polar/ionic solutes dissolve in polar solvents (like water) via ion-dipole or dipole-dipole/H-bonding interactions, while nonpolar solutes dissolve in nonpolar solvents via dispersion forces. Dissolution is favorable when the energy released forming new solute-solvent attractions offsets the energy needed to separate solute-solute and solvent-solvent particles, and when the increase in entropy (disorder) upon mixing is favorable. Most solid solutes become more soluble in water as temperature increases (endothermic dissolution favored by heat), whereas gas solubility in water decreases as temperature rises (gases escape faster when warmed); Henry's law states that gas solubility increases proportionally with the partial pressure of that gas above the liquid, which is why sealed carbonated beverages stay fizzy under pressure.

• •Polar/ionic solutes dissolve in polar solvents; nonpolar solutes dissolve in nonpolar solvents
• •Dissolution favorable when solute-solvent attraction compensates for separating solute and solvent particles
• •Solid solubility in water: generally increases with temperature
• •Gas solubility in water: decreases with temperature, increases with partial pressure (Henry's law)

Rank CH₄, CH₃OH, and CH₃F in order of increasing boiling point, and justify using IMFs.

1. 11. Identify strongest IMF in each: CH₄ is nonpolar → dispersion only.
2. 22. CH₃F has a polar C–F bond and no O-H/N-H/F-H bond on F attached to H → dipole-dipole (F is not bonded to H here).
3. 33. CH₃OH has an O–H bond → hydrogen bonding, the strongest of the three.
4. 44. Strength order: dispersion (CH₄) < dipole-dipole (CH₃F) < hydrogen bonding (CH₃OH).
5. 55. Boiling point follows IMF strength directly since molar masses are similar.

Increasing boiling point: CH₄ < CH₃F < CH₃OH

Calculate the total heat required to convert 25.0 g of ice at −10.0 °C to steam at 110.0 °C. (C_ice = 2.09 J/g°C, C_water = 4.18 J/g°C, C_steam = 2.01 J/g°C, ΔH_fus = 334 J/g, ΔH_vap = 2260 J/g)

1. 11. Heat ice from −10 to 0°C: q₁ = 25.0 g × 2.09 J/g°C × 10.0°C = 522.5 J
2. 22. Melt ice at 0°C: q₂ = 25.0 g × 334 J/g = 8350 J
3. 33. Heat water from 0 to 100°C: q₃ = 25.0 g × 4.18 J/g°C × 100.0°C = 10 450 J
4. 44. Vaporize water at 100°C: q₄ = 25.0 g × 2260 J/g = 56 500 J
5. 55. Heat steam from 100 to 110°C: q₅ = 25.0 g × 2.01 J/g°C × 10.0°C = 502.5 J
6. 66. Sum all steps: q_total = 522.5 + 8350 + 10450 + 56500 + 502.5 = 76 325 J

q_total ≈ 7.63 × 10⁴ J (76.3 kJ)

Silver crystallizes in a face-centered cubic (FCC) unit cell. How many silver atoms are contained per unit cell, and what is the coordination number?

1. 11. FCC has atoms at the 8 corners and the centers of the 6 faces.
2. 22. Corner atoms contribute 8 × 1/8 = 1 atom.
3. 33. Face atoms contribute 6 × 1/2 = 3 atoms.
4. 44. Total atoms per cell = 1 + 3 = 4 atoms.
5. 55. FCC has a coordination number of 12 (each atom touches 12 nearest neighbors).

4 Ag atoms per unit cell; coordination number = 12

Explain, using the phase diagram of water, why increasing pressure on ice at 0°C causes it to melt, while increasing pressure on most other solids causes them to freeze further.

1. 11. Water's solid-liquid boundary line has a negative slope (leans left) because ice is less dense than liquid water.
2. 22. At constant temperature, moving to higher pressure on a diagram means moving straight upward.
3. 33. For water, moving up at 0°C crosses the negatively-sloped line from the solid region into the liquid region — so ice melts.
4. 44. For a typical substance, the solid-liquid line slopes right (solid denser than liquid), so moving up at the melting point stays in/enters the solid region — favoring freezing.
5. 55. This is why ice skates and glaciers can locally melt ice under pressure, an anomaly caused by water's unique density behavior.

Water's ice is less dense than its liquid, giving a negatively sloped solid-liquid line, so pressure favors melting instead of freezing (unlike most substances).

A compound has a vapor pressure of 1 atm at 78°C and is used in a container at high altitude where atmospheric pressure is 0.80 atm. Will the compound's boiling point at this altitude be higher or lower than 78°C, and why?

1. 11. Boiling occurs when vapor pressure equals external (atmospheric) pressure.
2. 22. The normal boiling point (78°C) corresponds to vapor pressure = 1 atm exactly.
3. 33. At high altitude, atmospheric pressure is lower (0.80 atm), so the liquid needs less vapor pressure to boil.
4. 44. Vapor pressure reaches 0.80 atm at a lower temperature than it reaches 1 atm.
5. 55. Therefore the compound will boil at a temperature below 78°C at this altitude.

Lower than 78°C — reduced atmospheric pressure lowers the temperature needed to reach the boiling vapor pressure.

Q1. Which type of intermolecular force is present in ALL molecules, polar or nonpolar?

Q2. Hydrogen bonding requires hydrogen to be covalently bonded directly to which elements?

Q3. Why does n-pentane have a higher boiling point than its branched isomer neopentane, despite identical molar mass?

Q4. What is the defining condition for dynamic equilibrium between a liquid and its vapor in a closed container?

Q5. A liquid's normal boiling point is defined as the temperature at which vapor pressure equals:

Q6. On a heating curve, what is happening to the substance during a flat plateau region?

Q7. Which best explains water's unusually high boiling point compared to H₂S?

Q8. In a phase diagram, the triple point represents:

Q9. Why does water's solid-liquid line on a phase diagram slope to the left (negative slope)?

Q10. Which type of solid is characterized by cations surrounded by a 'sea' of delocalized electrons, giving malleability and electrical conductivity?

Q11. How many atoms are contained per unit cell in a body-centered cubic (BCC) structure?

Q12. According to 'like dissolves like,' which solvent would best dissolve nonpolar iodine (I₂)?

Q13. As temperature increases, the solubility of most gases in water:

Q14. A concave meniscus (curving upward at the edges) forms in a glass tube when:

Q15. Beyond the critical point on a phase diagram, a substance exists as: